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PHOSPHORIC ACID, META / PHOSPHORIC ACID, ORTHO
697
PHOSPHORIC ACID, META
[37267-86-0]
Formula HPO3; MW 79.98; the general formula (HPO3)n; in vapor phase it
probably exists as monomeric HPO3
Synonyms: metaphosphoric acid; glacial phosphoric acid. Long chain linear or
cyclic metaphosphoric acids of formula (HPO3)n are also called polymetaphosporic acids.
Uses
Metaphosphoric acid is used in making oxyphosphate dental cement. It also
is used as an analytical reagent.
Physical Properties
Glass-like colorless solid; soft and transparent; deliquesces; density 2.2 to
2.5 g/cm3; sublimes at red heat; dissolves slowly in cold water, decomposing to
phosphoric acid; soluble in alcohol.
Thermochemical Properties
∆Hƒ°
–226.7 kcal/mol
Preparation
Metaphosphoric acid is obtained as a polymeric glassy solid by prolonged
heating of phosphoric acid. Either phosphoric acid, H3PO4, or pyrophosphoric
acid, H4P2O7, when heated above 300°C, on cooling yields the transparent
glassy mass of composition (HPO3)n.
Metaphosphoric acid also is obtained from partial hydration of phosphorus
pentoxide by dissolving it in cold water.
Analysis
Elemental composition: P 38.73%, H 1.26%, O 60.01%. The compound may
be identified by physical properties alone. It may be distinguished from ortho
and pyrophosphates by its reaction with a neutral silver nitrate solution.
Metaphosphate forms a white crystalline precipitate with AgNO3, while PO43¯
produces a yellow precipitate and P2O73¯ yields a white gelatinous precipitate.
Alternatively, metaphosphate solution acidified with acetic acid forms a white
precipitate when treated with a solution of albumen. The other two phosphate
ions do not respond to this test. A cold dilute aqueous solution may be analyzed for HPO3̄ by ion chromatography using a styrene divinylbenzene-based
low-capacity anion-exchange resin.
PHOSPHORIC ACID, ORTHO
[7664-38-2]
Formula H3PO4; MW 97.995; also forms a hemihydrate H3PO4•1/2 H2O
[16271-20-8], known as diphosphoric acid; crystals of anhydrous acid or hemi-
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PHOSPHORIC ACID, ORTHO
hydrate consist of tetrahedral PO4 units linked by hydrogen bonding.
Structure (OH)3P=O.
Synonyms: phosphoric acid; trihydrogen phosphate.
History and Uses
Phosphoric acid was prepared first by Robert Boyle in 1694 by dissolving
phosphorus pentoxide in water. Phosphoric acid is probably the most important compound of phosphorus. It is the second largest inorganic chemical by
volume, after sulfuric acid, marketed in the United States. The single most
important application of this acid is manufacturing phosphate salts for fertilizers. Such fertilizer phosphates include sodium, calcium, ammonium, and
potassium phosphates. Other applications are in metal pickling and surface
treatment for removal of metal oxides from metal surfaces; electropolishing of
aluminum; as a bonding agent in various refractory products such as alumina and magnesia; as a catalyst in making nylon and gasoline; as a dehydrating agent; in fireproofing wood and fabrics; in lithographic engraving; in textile dyeing; in dental cement; in coagulating rubber latex; in purifying hydrogen peroxide; and as a laboratory reagent. Dilute solutions of phosphoric acid
are used as additives to carbonated beverages for a pleasing sour taste. Also,
dilute acid is used in refining sugar; as a nutrient; and as a buffering agent in
preparing jam, jelly, and antibiotics. The commercial phosphoric acid is 85%
(w/w) in strength.
Physical Properties
White orthorhombic crystals in pure and anhydrous state or a clear, syrupy
liquid; melts at 42.35°C; hygroscopic; can be supercooled into a glass-like
solid; crystallizes to hemihydrate, H3PO4•1/2H2O on prolonged cooling of 88%
solution; hemihydrate melts at 29.32°C and loses water at 150°C; density
1.834 g/cm3 at 18°C; density of commercial H3PO4 (85%) 1.685 g/mL at 25°C;
pH of 0.1N aqueous solution 1.5; extremely soluble in water, 548 g/100mL at
room temperature; soluble in alcohol.
Thermochemical Properties
∆H ƒ°
∆Gƒ°
S°
Cρ
∆H soln
–305.7 kcal/mol
–243.5 kcal/mol
26.41 cal/deg mol
25.35 cal/ degree mol
2.79 kcal/mol
Production
Low-purity technical grade phosphoric acid for use in fertilizers is produced from phosphate rocks by digestion with concentrated sulfuric acid. The
apatite types, primarily consisting of calcium phosphate phosphate rocks, are
used:
Ca3(PO4)2 + 3H2SO4 + 6H2O → 2H3PO4 + 3(CaSO4•2H2O)
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PHOSPHORIC ACID, ORTHO
699
The insoluble calcium sulfate slurry is filtered out. Acid from this wet
process is impure but can be purified by various methods. Purification steps
involve precipitation, solvent extraction, crystallization, and ion exchange
techniques.
Phosphoric acid also can be made by many different methods. Dissolution
of phosphorus pentoxide in water and boiling yields phosphoric acid. Pure
phosphoric acid can be obtained by burning phosphorus in a mixture of air
and steam:
P4 (l) + 5O2 (g) →P4O10 (s)
P4O10 (s) + H2O (g) → 4H3PO4 (l)
The acid also may be prepared by heating violet phosphorus with 33%
nitric acid:
4P + 10HNO3 + H2O → 4H3PO4 + 5NO ↑ + 5NO2 ↑
or by heating red phosphorus with nitric acid (1:1). The overall equation is:
P + 3HNO3 → H3PO4 + NO + 2NO2
Reactions
Phosphoric acid is a tribasic acid. It is not an oxidizing acid. In aqueous
solution phosphoric acid dissociates to H2PO4̄ , HPO42¯ and PO33¯ ions. The
dissociation constants are as follows:
H3PO4 + H2O ↔ H3O+ + H2PO4̄
Ka1 = 7.1×10–3
H2PO4̄ + H2O ↔ H3O+ + HPO42–
Ka2 = 8.0×10–8
HPO42– +H2O ↔ H3O+ + PO43–
Ka3 4.8×10–13
Thus, out of the three ionizable hydrogens in phosphoric acid, the first H+ is
removed more easily than the second, and the second H+ dissociates more easily than the third. When phosphoric acid is titrated with sodium hydroxide,it
forms both acidic and basic salts:
H3PO4 + NaOH → NaH2PO4 + H2O
(moderately acidic)
H3PO4 + 2NaOH → Na2HPO4 + 2H2O
(very weakly basic)
H3PO4 + 3NaOH → Na3PO4 + 3H2O
(weakly basic)
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PHOSPHORIC ACID, ORTHO
At first, the pH of the solution suddenly changes to 4.5 upon completing formation of NaH2PO4, and there is a second change of pH to around 9.0 after
Na2HPO4 is formed completely. The titration curve for H3PO4 =NaOH shows
two pH inflection points. Mixtures of Na2HPO4 and NaH2PO4, therefore,
exhibit buffer properties at pH 6 to 8. Because of a big difference in the first
two ionization constants, phosphoric acid can be titrated as a monobasic or a
dibasic acid. Aqueous solution of trisodium phosphate, Na3PO4 is basic.
Heating phosphoric acid converts it to metaphosphoric acid, HPO3, and
pyrophosphoric acid, H4P2O7. If anhydrous phosphoric acid is melted and
allowed to stand for several weeks, it partially converts to pyrophosphoric
acid, H4P2O7. The equilibrium occurs in liquid phase:
2H3PO4 ↔ H4P2O7 + H2O
However, when heated at 215°C, it is fully converted to pyrophosphoric acid.
When heated at 300°, phosphoric acid converts to metaphosphoric acid,
HPO3:
H3PO4 → HPO3 + H2O
Phosphoric acid reacts with most metals and their oxides at temperatures
above 400°C forming metal phosphates. Reactive metals, such as magnesium,
react with phosphoric acid solutions forming magnesium phosphate and
evolving hydrogen:
2H3PO4 + 3Mg → Mg3(PO4)2 + 3H2
When ammonia gas is bubbled through phosphoric acid solution, diammonium hydrogen phosphate is produced:
H3PO4 (aq) + 2NH3 (g) → (NH4)2HPO4 (s)
Reaction with calcium triphosphate fluoride yields calcium dihydrogen
phosphate, a component of superphosphate fertilizer:
7H3PO4 + Ca(PO4)3F + 5H2O → 5 Ca(H2PO4)HPO4 + HF
When heated with solid sodium bromide, phosphoric acid yields sodium
dihydrogen phosphate, liberating hydrogen bromide:
H3PO4 (l) + NaBr (s) → NaH2PO4 (s) + HBr (g)
Analysis
The orthophosphate anion, PO43¯ can be analyzed readily in aqueous solution by either ion chromatography or colorimetry. The aqueous solution must
be diluted for such analyses. In colorimetric measurement, the solution is
treated with a reagent mixture containing ammonium molybdate and ammo-
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PHOSPHORIC ACID, PYRO
701
nium metavanadate under acid conditions to form the yellow color of vanadomolybdophosphoric acid. The yellow absorbance or transmittance may be
measured at λ 470nm. Alternatively, the solution after treatment with ammonium molybdate may be reduced by stannous chloride to produce an intense
blue color of molybdenum blue which may be measured at λ 650 nm. The concentration of PO43¯ may be calculated from a standard calibration curve. The
normality of phosphoric acid can be measured by titration with a standard
solution of NaOH using a suitable color indicators or a pH meter.
PHOSPHORIC ACID, PYRO
[2466-09-3]
Formula: H4P2O7; MW 177.98
Structure: (HO)2P(=O)—O—P(=O)(OH)2
Synonym: disphosphoric acid
Uses
No commercial application of this acid is known. The pyrophosphate salts
usually are not made from the acid.
Physical Properties
Colorless needles or liquid; hygroscopic. Crystallizes in two anhydrous
forms: a metastable form melting at 54.3°C and a second and more stable form
melting at 71.5°C; extremely soluble in cold water, reacting very slowly to
form phosphoric acid; decomposing much faster in hot water; very soluble in
alcohol and ether.
Preparation
Pyrophosphoric acid may be prepared by heating orthophosphoric acid at
215°C:
2H3PO4 → H4P2O7 + H2O
The acid solution in pure form can be obtained by ion exchange, passing an
aqueous solution of sodium pyrophosphate, Na4P2O7, through a suitable
cation exchange column.
Reactions
The acid has four replacable H+ ions. Its dissociation constants indicate
that two H+ ions are strongly acidic while the other two protons are weakly
acidic. The first dissociation constant especially is very large:
H4P2O7 + H2O ↔ H3O+ + H3P2O7̄
Ka1 ~ 10–1
H3P2O7̄ + H2O ↔ H3O+ + H2P2O72–
Ka2 ~ 1.5x10–2
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702 PHOSPHORUS
H2P2O72¯ + H2O ↔ H3O+ + HP2O73–
Ka3 ~ 2.7x10–7
HP2O73¯ + H2O ↔ H3O+ + P2O74¯
Ka4 ~ 2.4x10–10
Pyrophosphoric acid forms acid salts, such as NaH3P2O7 and Na2H2P2O7.
Analysis
Pyrophsophoric acid may be converted into its acid salts which may be
characterized individually by physical and x-ray properties and elemental
compositions.
PHOSPHORUS
[7723-14-0]
Symbol P; atomic number 15; atomic weight 30.974; a Group VA (Group 15)
nonmetallic element of nitrogen group; electron configuration [Ne]3s23ρ3;
valence states –3, +3, +5; most stable valence state +3; atomic radius 1.10Å;
one natural isotope P-31 (100%); nine radioactive isotopes in the mass range
26, 28–30, 32–36; the longest-lived radioisotopes P-33, t ½ 25.3 day
History, Occurrence, and Uses
Elemental phosphorus was discovered in 1669 by Hennig Brand. About two
hundred years later James Readman developed a process for phosphorus
recovery from phosphatic rocks using an electric furnace.
Phosphorus is one of the most widely distributed elements on earth. It is
found as phosphate salts in nearly all igneous rocks and in sedimentary
deposits and sea beds. Phosphorus occurs in more than three hundred minerals, usually associated with Ca, Mg, Fe, Sr, Al, Na, and several other metals,
and with anions such as silicates, sulfates, oxides, hydroxides, and halides.
Phosphorus is an essential element present in all living matter and is vital
in biological and ecological processes. It occurs in DNA and other nucleic
acids, and in bones.
Phosphorus is used in pyrotechnics, smoke bombs, incendiary shells, and
safety matches. It also is used in organic syntheses, manufacture of phosphoric acid, phosphorus trichloride, phosphine, and other compounds.
Physical Properties
Elemental phosphorus in solid phase exists in three major allotropic forms:
(1)white or yellow phosphorus that may occur in alpha or beta modification,
(2) red phosphorus, and (3) black phosphorus.
White phosphorus is a white, soft, wax-like transparent mass which often
acquires a yellow appearance due to impurities, especially traces of red phosphorus. It has a garlic-like odor. It is made up of cubic crystals, has a density
1.82 g/cm3, and melts at 44.1°C to a colorless or yellowish liquid. X-ray diffraction studies and 31P-NMR analysis indicate tetrahedral P4 molecules with
an interatomic distance of 2.21Å , and the molecules are able to rotate freely
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703
in the crystals. When cooled below –76.9°C, the cubic alpha form converts to
a hexagonal beta modification with a density 1.88 g/cm3. The beta form,
unlike the alpha form, does not rotate freely in the crystal but has a fixed orientation of P4 molecules in the lattice.
Red phosphorus is obtained from white phosphorus by heating at 230 to
240°C, allowing complete conversion to occur in about 48 hours. Conversion is
catalyzed by sulfur, iodine, and selenium. The red allotrope also slowly
deposits from liquid phosphorus or from a solution of white phosphorus, the
rate and yield depending on catalysts, temperature, light, and other factors.
Red phosphorus exhibits various modifications. Three important ones are an
amorphous form at ordinary temperatures and two crystalline modifications
which include a triclinic form and a hexagonal or a tetragonal form that may
prevail at higher temperatures. There also are a few more modifications, all
of which may coexist, accounting for variability in physical properties of red
phosphorus. The triclinic variety of red phosphorus is the most stable of all
allotropes of phosphorus at ordinary temperatures. Red phosphorus possesses a density of 2.0 to 2.31 g/cm3 and melts at 590°C.
Black phosphorus is the third major allotropic form of phosphorus. It
occurs in two forms, one is an amorphous modification having a laminar structure similar to graphite and the other is an orthorhombic crystalline form.
The density of black phosphorus may vary between 2.20 to 2.69 g/cm3. Black
phosphorus is obtained from white phosphorus by heating the latter at 220°C
under an extremely high pressure of about 10,000 atm.
When solid phosphorus of any form—white, red, or black—is melted, it
forms the same liquid phosphorus. This liquid has a density of 1.74 g/cm3 and
viscosity 1.69 centipoise at 50°C. Liquid phosphorus boils at 280.5°C. Upon
cooling, liquid phosphorus solidifies to only white phosphorus. Liquid phosphorus and its vapors consist of tetrahedral P4 molecules. The vapors, on
rapid condensation, convert to white phosphorus.
While white and red phosphorus have high electrical resistivity, the black
variety has a low resistivity of 0.71 ohm-cm at 0°C. Solubility also varies widely. White phosphorus is soluble in a number of organic solvents. It is very
highly soluble in carbon disulfide, about 400 g/100 g solvent at 0°C and moderately soluble in benzene (~3.59 g/100g at 25°C) and exhibits lower solubility in ether (~1.5g/100g at 25°C). Red and black phosphorus are insoluble in
organic solvents. White phosphorus is a flammable solid, igniting spontaneously in air at 35°C. Red and black phosphorus are nonflammable. The latter is difficult to ignite.
Production
White phosphorus usually is obtained by heating some form of calcium
phosphate with quartz and coke, usually in an electric furnace. The reactions
may be written in two steps as follows:
Ca3(PO4)2 + 3SiO2 → 3CaSiO3 + P2O5
P2O5 + 5C → 2P + 5 CO
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PHOSPHORUS
In commercial scale, white phosphorus is manufactured mostly from the
mineral fluorapatite by heating with silica and coke in an electric-arc or blast
furnace at a temperature of 1,200 to 1,500°C. An overall reaction may be represented in the following equation.
4Ca5F(PO4)3 + 18SiO2 + 30C → 18CaO • SiO2 • 2CaF2 + 30CO↑ + 3P4↑
(slag)
White phosphorus also can be produced by a wet process using phosphoric
acid, a process that was practiced historically in commercial production. In
this method the starting material, phosphoric acid, usually is prepared in
large vats by reacting phosphate rock with sulfuric acid:
Ca5F(PO4)3 + 5H2SO4 + 10H2O → 3H3PO4 + 5CaSO4 • 10H2O + HF
Phosphoric acid is filtered out of the mixture. It is then mixed with coke,
charcoal or sawdust; dried; charred; and finally heated to white heat in a fireclay retort:
H3PO4 + 16C → P4 + 6H2 + 16CO
The vapor is condensed to obtain white phosphorus.
As stated earlier, all other forms of phosphorus can be made from white
phosphorus. Thus, heating white phosphorus first at 260°C for a few hours
and then at 350°C gives red phosphorus. The conversion is exothermic and
can become explosive in the presence of iodine as a catalyst. When a solution
of white phosphorus in carbon disulfide or phosphorus tribromide is irradiated the scarlet red variety is obtained.
Black phosphorus allotrope is produced by heating white phosphorus at
220°C under 12,000 atm pressure. The conversion is initially slow, but can
became fast and explosive after an induction period.
White phosphorus is stored under water as it ignites in air. It may be cut
into appropriate sizes only under water.
Reactions
Reactivity of white phosphorus is much greater than red or black phosphorus. Black phosphorus is the least reactive of all phosphorus allotropes.
White phosphorus ignites in air spontaneously. When placed on a paper,
the paper catches fire after a short delay. It catches fire at about 35°C. At
room temperature white phosphorus glows in the dark on exposure to air
emitting faint green light. Such chemiluminescence is attributed to the oxidation of P4 molecules in the vapor phase in contact with the surface of solid
phosphorus:
P4(g) + 5O2(g) → P4O10(s) + light
The mechanism involves a complicated oxidative process that occurs only at
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PHOSPHORUS
705
certain partial pressures of oxygen and not in pure oxygen at atmospheric
pressure, nor in vacuum.
Red phosphorus ignites when struck with a hammer blow or when
heated at 260°C. Black phosphorus ignites in contact with flame.
White phosphorus reacts spontaneously with halogens at ordinary
temperatures forming phosphorus trihalides. However, in excess halogen the
product is phosphorus pentahalide:
P4(s) + 6Cl2(g) → 4PCl3 (l)
P4 (s) + 10Cl2 (g) → 4PCl5 (s)
White phosphorus reacts with sulfur on warming forming phosphorus
trisulfide:
P4(s) + 6S(s) → 2P2S3 (s)
White phosphorus reacts with strong aqueous alkali solution forming
hypophosphite with evolution of phosphine, PH3:
P4 + 3KOH + 3H2O → 3KH2PO2 + PH3 ↑
Strong oxidzing agents, such as nitric acid (cold and concentrated), oxidize
phosphorus to phosphoric acid.
Reaction with copper sulfate solution forms a mixture of metallic copper
and copper(I) phosphide. The two reactions may be written separately as follows:
P4 + 10CuSO4 + 16H2O → 10Cu +4H3PO4 + 10H2SO4
3P4 + 12CuSO4 + 24H2O → 4Cu3P + 8H3PO3 + 12H2SO4
Similar reactions occur with the salts of other easily reducible metals,
such as silver and gold, in aqueous salt solutions.
Phosphorus combines with several metals on heating, forming their phosphides.
P4 + 6Ca → 2Ca3P2
Reactions with alkali metals occur under warm conditions producing the
corresponding metal phosphides:
P4 + 12Na → 4Na3P
Analysis
The allotropes of phosphorus may be identified from their physical properties. White phosphorus can be identified from its chemiluminescence (a pale
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PHOSPHORUS ACID
green glow) at a specific range of oxygen partial pressure at room temperature. Furthermore, it spontaneously ignites in air at 35°C. It also imparts
chemiluminescence to water when boiled. Elemental phosphorus can be analyzed by GC/MS. Its solution in a suitable organic solvent, such as benzene
may be injected, onto the GC and identified from the mass spectra. In solution
it exists as P4 molecule, thus the characteristic molecular ion should have the
mass 124. Red phosphorus can be converted into its white allotrope by heating in the absence of air to above 260°C and condensing the vapors and trapping in an organic solvent for analysis by GC/MS.
Hazards
White phosphorus is a highly toxic substance, both an acute and chronic
toxicant. Chronic exposure to it’s vapors can cause “phossy jaw;” necrosis of
the jaw. Other symptoms are bronchopneumonia, bone changes, anemia and
weight loss, Ingestion can cause nausea, vomiting, abdominal pain, diarrhea
and coma. Skin contact can cause severe burns. In the eye it damages vision.
Red phosphorus is much less toxic than its white allotrope. Its fumes, when
burned, are highly irritating. White phosphorus is a flammable solid, igniting
spontaneously when exposed to air.
PHOSPHORUS ACID
[13598-36-2]
Formula: H3PO3; MW 82.00; a dibasic acid
Structure: HP(=O)(OH)2
Synonym: orthophosphorus acid
Uses
Phosphorus acid is used to prepare phosphite salts. It is usually sold as a
20% aqueous solution.
Physical Properties
White crystalline mass; deliquescent; garlic-like odor; density 1.651 g/cm3
at 21°C; melts at 73.6°C; decomposes at 200°C to phosphine and phosphoric
acid; soluble in water, about 310 g/100mL; K1 5.1x10–2 and K2 1.8x10–7; soluble in alcohol.
Thermochemical Properties
∆Hƒ°
–230.5 kcal/mol
Preparation
Phosphorus acid can be prepared by the reaction of phosphorus
trichloride with water:
PCl3 + 3H2O → H3PO4 + 3HCl
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PHOSPHORUS ACID
707
The reaction is violent. Addition of PCl3 should be extremely cautious and
slow. The addition can be carried out safely in the presence of concentrated
HCl. Alternatively, a stream of air containing PCl3 vapor is passed into icecold water and solid crystals of H3PO4 form.
Alternatively, phosphorus acid can be prepared by adding phosphorus
trichloride to anhydrous oxalic acid:
PCl3 + 3(COOH)2 → H3PO3 + 3CO + 3CO2 + 3HCl
In this reaction, all products except H3PO3 escape as gases leaving the liquid acid.
Dissolution of phosphorus sesquioxide in water also forms phosphorus acid.
When shaken with ice water, phosphorus acid is the only product
.
P4O6 + 6H2O → 4H3PO3
However, in hot water part of the phosphorus acid disproportionates to
phosphoric acid and phosphorus or phosphine.
Reactions
Phosphorus acid on heating at 200°C converts to phosphoric acid and phosphine:
4H3PO3 → 3H3PO4 + PH3
Phosphorus acid is a moderately strong dibasic acid. It reacts with alkalies
forming acid phosphites and normal phosphites. Thus, reaction with sodium
hydroxide gives sodium dihydrogen phosphite and disodium hydrogen phosphite, but not sodium phosphate, Na3PO4.
H3PO3 + NaOH → NaH2PO3 + H2O
H3PO3 + 2NaOH → Na2HPO3 + 2H2O
Phosphorus acid is a powerful reducing agent. When treated with a cold
solution of mercuric chloride, a white precipitate of mercurous chloride forms:
H3PO3 + 2HgCl2 + H2O → Hg2Cl2 + H3PO4 + 2HCl
Mercurous chloride is reduced further by phosphorus acid to mercury on
heating or on standing:
H3PO3 + Hg2Cl2 + H2O → 2Hg + H3PO4 + 2HCl
Phosphorus acid reacts with silver nitrate in dilute solution yielding a
white precipitate of silver phosphite, Ag3PO3 , which reduces to metallic silver.
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PHOSPHORUS OXYCHLORIDE
H3PO3 + 2Ag3PO3 + H2O → 6Ag + H3PO4 + 2HPO3
Analysis
Elemental composition: P 37.78%, H 3.69%, O 58.54%. The acid in solid form
may be identified by its physical properties. Aqueous solution may be heated
and phosphorus acid is converted to phosphoric acid which is measured for
orthophosphate ion by ion chromatography or colorimetry (see Phosphoric
Acid). A cold aqueous solution may be analyzed for phosphite ion by ion chromatography, following appropriate dilution. Strength of the acid in an aqueous solution may be measured by acid-base titration using a standard solution
of alkali. Also, titration against a standard solution of silver nitrate using
potassium chromate as indicator may serve as an additional confirmatory
test.
PHOSPHORUS OXYCHLORIDE
[100025-87-3]
Formula: POCl3; MW 153.33
Synonym: phosphoryl chloride
Uses
Phosphorus oxychloride is a chlorinating agent in many organic preparative reactions. It also is a solvent in cryoscopy.
Physical Properties
Colorless fuming liquid with a pungent odor; density 1.645 g/mL; freezes at
1°C; boils at 105.5°C; reacts with water and ethanol.
Thermochemical Properties
∆Ηƒ° (liq)
∆Ηƒ° (gas)
∆Gƒ° (liq)
∆Gƒ° (gas)
S° (liq)
S° (gas)
Cρ (liq)
Cρ (gas)
∆Hfus
∆Hvap
–142.7 kcal/mol
–133.5 kcal/mol
124.5 kcal/mol
–122.6 kcal/mol
53.2 cal/deg mol
77.8 cal/deg mol
33.2 cal/deg mol
20.3 cal/deg mol
3.13 kcal/mol
8.21 kcal/mol
Preparation
Phosphorus oxychloride can be prepared from phosphorus trichloride or
phosphorus pentachloride. It can be obtained from phosphorus trichloride by
cautious addition of potassium chlorate:
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PHOSPHORUS PENTACHLORIDE
709
3PCl3 + KClO3 → 3POCl3 + KCl
The oxychloride also is obtained by the action of boric acid or oxalic acid
with phosphorus pentachloride:
3PCl5 + 2B(OH)3 → 3POCl3 + B2O3 + 6HCl
PCl5 + (COOH)2 → POCl3 + CO + CO2 + 2HCl
Phosphorus oxychloride also is made by heating calcium phosphate in a
current of chlorine and carbon monoxide at 350°C:
2Ca3(PO4)2 + 9Cl2 + 6CO → 4POCl3 + 6CaCO3
Alternatively, heating a mixture of calcium phosphate and carbon in a current of chlorine at 750°C yields the oxychloride.
Reactions
Phosphorus oxychloride hydrolyzes in water forming phosphoric acid:
POCl3 +3H2O → H3PO4 + 6HCl
When the vapors of phosphorus oxychloride are passed over carbon at red
heat, phosphorus trichloride is produced:
POCl3 + C → PCl3 + CO
The oxychloride also is reduced by hydrogen, carbon monoxide and other
reducing agents.
Analysis
Elemental composition: P 20.20%, O 10.43%, Cl 69.36%. The compound is
hydrolyzed in water and the products phosphoric and hydrochloric acids are
measured by a colorimetric method for orthophosphate ion (see Phosphoric
Acid, Analysis), and titration with silver nitrate for the chloride ion. Also,
phosphate and chloride ions can be measured by ion chromatography.
Toxicity
The compound is highly irritating to skin, eyes and mucous membranes.
Inhaling its vapors can cause pulmonary edema.
PHOSPHORUS PENTACHLORIDE
[10026-13-8]
Formula: PCl5; MW 208.24
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PHOSPHORUS PENTACHLORIDE
Uses
Phosphorus pentachloride is used as a chlorinating agent in many organic
syntheses, such as production of alkyl and acid chlorides. It also is a catalyst
in manufacturing acetylcellulose.
Physical Properties
Yellowish-white tetragonal crystals; pungent odor; fumes in air; deliquescent; density 2.1 g/cm3; decomposes on heating; melts at 166.8°C under the
pressure of its own vapor(triple point); sublimes at 160°C; critical temperature 373°C; hydrolyzes in water; soluble in carbon disulfide and carbon tetrachloride.
Thermochemical Properties
∆Hƒ°
–106 kcal/mol
Preparation
Phosphorus pentachloride is prepared by reacting white phosphorus with
excess dry chlorine. The white phosphorus is placed over sand in a retort from
which air and moisture have been purged. The reaction is indicated by inflaming phosphorus:
P4 + 10Cl2 → 4PCl5
Also, the compound is obtained by reaction of dry chlorine with phosphorus
trichloride:
PCl3 + Cl2 → PCl5
Reactions
Phosphorus pentachloride absorbs moisture from air forming phosphoryl
chloride:
PCl5 + H2O → POCl3 + 2HCl
The above reaction is difficult to control and progresses to complete hydrolysis. Thus, in the presence of excess water or when treated with water, the
pentachloride is hydrolyzed to phosphoric acid:
PCl5 + 4H2O → H3PO4 + 5HCl
Reaction with sulfur dioxide yields thionyl chloride and phosphoryl chloride:
PCl5 + SO2 → SOCl2 + POCl3
Reaction with liquid hydrogen sulfide forms thiophosphoryl chloride,
PSCl3:
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PHOSPHORUS PENTAFLUORIDE
711
PCl5 + H2S → PSCl3 + 2HCl
Phosphorus pentachoride converts arsenic to arsenic trichloride:
3PCl5 + 2As → 3AsCl3 + 3PCl3
Reaction with oxalic acid or boric acid yields phosphoryl chloride:
PCl5 + (COOH)2 → POCl3 + CO + CO2 + 2HCl
3PCl5 + 2B(OH)3 → 3POCl3 + B2O3 + 6HCl
Reaction with phosphorus pentoxide produces phosphoryl chloride:
3PCl5 + P2O5 → 5POCl3
Analysis
Elemental composition: P 14.88%, Cl 85.12%. The compound may be
hydrolyzed with water and the products phosphoric and hydrochloric acids
are measured for phosphate and chloride ions by ion chromatography and colorimetric methods (see Phosphoric Acid, Hydrochloric Acid).
Toxicity
The compound is strongly irritating to skin, eyes and mucous membranes.
PHOSPHORUS PENTAFLUORIDE
[7647-19-0]
Formula: PF5; MW 125.97
Synonym: phosphorus(V)fluoride
Uses
Phosphorus pentafluoride is a catalyst in ionic polymerization reactions.
Physical Properties
Colorless gas; fumes in air; density 5.527g/L; heavier than air, density in
air 4.35 (air=1); liquefies at –84.6°C; freezes at –93.8°C; reacts with water.
Thermochemical Properties
∆Ηƒ°
–381.1 kcal/mol
∆Gƒ°
–363.5 kcal/
S°
1.9 cal/deg mol
Cρ
20.3 cal/ deg mol
∆Hap
4.11 kcal/mol
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PHOSPHORUS PENTAFLUORIDE
Preparation
Phosphorus pentafluoride may be prepared by several methods, among
which are:
1. Treating phosphorus trifluoride with bromine and then heating the product phosphorus trifluoride dibromide, PF3Br2:
PF3 + Br2 → PF3Br2
5PF3Br2 → 3PF5 + 2PBr5
2. Heating phosphorus pentachloride with arsenic trifluoride:
PCl5 + 5AsF3 → 3PF5 + 5AsCl3
3. Subjecting phosphorus trifluoride to an electric spark in the absence of
air (a disproportion reaction occurs):
5PF3 → 3PF5 + 2P (in the presence of air, the product is phosphorus oxyfluoride, POF3)
4. Heating a mixture of phosphorus pentoxide and calcium fluoride:
P2O5 + 5CaF2 → 2PF5 + 5CaO
5. Heating a mixture of phosphorus oxyfluoride, hydrogen fluoride and sulfur trioxide:
POF3 + 2HF + SO3 → PF5 + H2SO4
The gas should be stored in steel cylinders in the absence of moisture.
Reactions
Phosphorus pentafluoride hydrolyzes in water, the products formed depend
on the reaction conditions. When exposed to moisture it forms phosphorus
oxyfluoride:
PF5 + H2O → POF3 + 2HF
Hydrolysis with water proceeds through formation of intermediates, oxyfluophosphates and ultimately gives phosphoric acids. The overall reaction may
be written as follows:
PF5 + 4H2O → H3PO4 + 5HF
The pentafluoride also is known to form adducts. With nitrogen dioxide it
forms an adduct, PF5 • NO2, at –10°C which decomposes on warming.
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PHOSPHORUS PENTOXIDE
713
Many complexes also are known, particularly with amines, pyridine, sulfoxides, and ethers.
Analysis
Elemental composition: P 24.59%, F 75.41%. The compound may be completely hydrolyzed in water and the ultimate hydrolysis products, phosphoric
acid and HF, may be determined for PO43¯and F¯ by ion chromatography. The
compound may be confirmed by GC/MS. Its diluent mixture in helium or other
inert gas may be introduced onto the GC column and PF5 may be identified
from its mass spectra. The characteristic mass ions are 126, 31, 107.
Toxicity
Phosphorus pentafluoride is a highly toxic gas. Inhalation can cause severe
irritation of mucous membrane and pulmonary edema. It is corrosive to skin
and can damage eyes.
PHOSPHORUS PENTOXIDE
[1314-56-3]
Formula: P2O5; MW 141.95; exists as P4O10 units as molecular entities
Synonyms: phosphorus pentaoxide; phosphorus(V) oxide; phosphoric anhydride
Uses
Phosphorus pentoxide is a very effective drying and dehydrating agent. It
also converts acids to their anhydrides.
Physical Properties
White, deliquescent, powdery solid; exhibits polymorphism; converts to several different crystalline forms on heating; the commercial material consists
of hexagonal crystals; the hexagonal crystals on very rapid heating first melt
at 420°C and then resolidify immediately to glassy orthorhombic crystals;
slow heating of hexagonal crystals causes melting at 340°C which, on solidification, gives the same metastable orthorhombic form; the glassy material
melts at about 580°C to a colorless and heavily viscous liquid; sublimes at
360°C; density of the commercial product 2.39g/cm3; reacts with water.
Thermochemical Properties
∆Hƒ° (hexagonal)
∆Hƒ° (amorphous)
∆Gƒ° (hexagonal)
S° (hexagonal)
Cρ (hexagonal)
–713.2 kcal/mol
–727.0 kcal/mol
–644.8 kcal/mol
54.7 cal/deg mol
50.6 cal/ deg mol
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PHOSPHORUS PENTOXIDE
Preparation
Phosphorus pentoxide is prepared by burning phosphorus in a plentiful
supply of dry air or oxygen:
P4 + 5O2 → P4O10
The crude product may contain a small amount of sesquioxide, P2O3, which
may be removed by sublimation in ozonized oxygen.
Reactions
The most important reaction of phosphorus pentoxide is its hydrolysis. The
hexagonal form reacts with water vigorously to form metaphosphoric acid
(HPO3)n which hydrolyzes further to yield phosphoric acid, H3PO4:
P4O10 + 2H2O → 4HPO3
HPO3 + H2O → H3PO4
In finely divided hexagonal form, the compound reacts violently with
water. The orthorhombic allotrope reacts much less vigorously than the
hexagonal form.
Phosphorus pentoxide dehydrates nitric acid at low temperatures (about
–10°C) forming metaphosphoric acid and nitrogen pentoxide:
P4O10 (s) + 4HNO3 (l) → 4HPO3( s) + 2N2O5 (s)
Reaction with phosphorus pentabromide yields phosphorus oxybromide:
P2O5 + 3PBr5 → 5POBr3
Analysis
Elemental composition: P 43.64%, O 56.36%. The pentoxide is dissolved in
water and the ultimate hydrolysis product, H3PO4, is analyzed for PO43– by
ion chromatography. Alternatively, the solution is treated with ammonium
molybdate—ammonium vanadate reagent to produce a yellow colored vanadomolybdophosphoric acid. Absorbance or transmittance of the solution may be
measured at a wavelength between 400 to 490 nm, depending on. concentration of PO43–. The solution must be diluted for analysis. The solution may further be reduced with stannous chloride to form an intensely colored molybdenum blue for measuring absorbance or transmittance at 690nm.
Toxicity
Phosphorus pentoxide is a strong irritant. It is corrosive to skin and contact
with eyes can be injurious.
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PHOSPHORUS TRICHLORIDE
715
PHOSPHORUS TRICHLORIDE
[7719-12-2]
Formula: PCl3; MW 137.33
Uses
Phosphorus trichloride is used to prepare phosphine and other phosphorus
compounds.
Physical Properties
Colorless fuming liquid; pungent odor; refractive index 1.516 at 14°C; density 1.574g/mL at 21°C; boils at 76°C; freezes at –112°C; decomposes in water;
soluble in benzene, carbon disulfide, ether and chloroform and other halogenated organic solvents.
Thermochemical Properties
∆Hƒ° (liq)
∆Hƒ° (gas)
∆Gƒ° (liq)
∆Gƒ° (gas)
S° (liq)
Cρ (liq)
–76.4 kcal/mol
–68.6 kcal/mol
–65.1 kcal/mol
–64.0 kcal/mol
74.5 cal/deg mol
17.17 cal/deg mol
Preparation
Phosphorus trichloride is prepared by reacting white phosphorus with dry
chlorine present in limited quantity. Excess chlorine will yield phosphorus
pentachloride, PCl5.
P4 + 6Cl2 → 4PCl3
P4 + 10Cl2 → 4PCl5
The compound is prepared in a retort attached to inlet tubes for dry chlorine and dry carbon dioxide and a distillation flask. White phosphorus is
placed on sand in the retort. All air, moisture, and any phosphorus oxide
vapors present in the apparatus are expelled by passing dry carbon dioxide.
Dry chlorine is then introduced into the apparatus. If a flame appears on
phosphorus it indicates presence of excess chlorine. In that event, the rate of
chlorine introduction should be decreased. For obtaining phosphorus trichloride, flame should appear at the end of the chlorine-entry tube. The trichloride formed is collected by condensation in the distillation flask. A soda lime
tube is attached to the apparatus to prevent moisture entering the flask.
Phosphorus trichloride also can be prepared by reducing phosphorus oxychloride vapors with carbon at red heat:
POCl3 + C → PCl3 + CO
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PHOSPHORUS TRICHLORIDE
Reactions
Phosphorus trichloride reacts violently with water forming phosphorus
acid:
PCl3 + 3H2O → H3PO3 + 3HCl
When dry chlorine gas is passed into the liquid trichloride under cooling,
phosphorus pentachloride is obtained:
PCl3 + Cl2 → PCl5
Phosphorus trichloride reacts with concentrated sulfuric acid forming
chlorosulfuric acid and metaphosphoric acid:
PCl3 + 2H2SO4 → HSO3Cl + HPO3 + SO2 + 2HCl
Reaction with sulfur trioxide produces phosphoryl chloride:
PCl3 + SO3 → POCl3 + SO2
Oxidation of trichloride produces phosphorus trichloride oxide which may
be used as the starting material to prepare alkyl-and alkoxyphosphine oxides:
PCl3 + O2 → O=PCl3
O=PCl3 + 3CH3OH → (CH3O)3P=O + 3HCl
O=PCl3
RMgX
→
R3P=O
Phosphorus trichloride reacts with thionyl chloride to form phosphoryl
chloride, thiophosphoryl chloride and phosphorus pentachloride:
3PCl3 + SOCl2 → POCl3 + PSCl3 + PCl5
It reacts violently with potassium chlorate forming phosphoryl chloride:
3PCl3 + KClO3 → 3POCl3 + KCl
Phosphorus trichloride reacts with iodine in warm glacial acetic acid solution, which on cooling yields orange crystals of phosphorus diiodide:
2PCl3 + 5I2 → P2I4 + 6ICl
Reaction with potassium iodide yields phosphorus triiodide:
PCl3 + 3KI → PI3 + 3KCl
Phosphorus trichloride reacts with organics that contain hydroxyl groups.
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PHOSPHORUS TRICHLORIDE
717
However, with ethanol two competing reactions occur:
PCl3 + 3C2H5OH → H3PO3 + 3C2H5Cl
PCl3 + 3C2H5OH →P(OC2H5)3 + 3HCl
With acetic acid the products are acetyl choride and phosphorus acid:
PCl3 + 3CH3COOH → 3CH3COCl + H3PO3
Similar reactions occur with other carboxylic acids.
Reactions with ammonia under controlled conditions produce phosphorus
triamine:
PCl3 + 3NH3 → P(NH2)3 + 3HCl + PCl3
Reaction with sulfur forms phosphorus trichloride sulfide:
PCl3 + S → S = PCl3
Phosphorus trichloride is converted to phosphorus trifluoride by heating
with flouride of arsenic, anitmony or zinc:
2PCl3 + 3ZnF2 → 2PF3 + 3ZnCl2
Reaction with silver isocyanate or silver thiocyanate yields phosphorus triisocyanate or phosphorus trithiocyanate:
PCl3 + 3AgNCO → P(NCO)3 + 3AgCl
PCl3 + 3AgSCN → P(SCN)3 + 3AgCl
Reaction with lower alcohols in the presence of a base yields the corresponding trialkoxyphosphine:
PCl3 + 3C2H5OH → P(OC2H5)3 + 3HCl
However, in the absence of a base the product is dialkoxyphosphine oxide,
(C2H5O)2PH(=O).
Phosphorus trichloride forms a tetracoordinated nickel complex by action
with nickel tetracarbonyl:
4PCl3 + Ni(CO)4 → P[Ni(PCl3)]4 + 4CO
Analysis
Phosphorus trichloride may be dissolved in a suitable organic solvent such
as benzene or chloroform and analyzed by GC-NPD in phosphorus mode. Its
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PLATINIC ACID, HEXACHLORO
solution in CS2 may be analyzed by GC—FID. The most definitive test is by
mass spectrometry.
Hazard
Phosphorus trichloride is highly corrosive. Its vapors are an irritant to
mucous membranes. Chronic exposure to its vapors can cause bronchitis. It
reacts violently with water and explodes in contact with acetic and nitric
acids, and several other substances (Patnaik. P. 1999. A Comprehensive Guide
to the Hazardous Properties of Chemical Substances, 2nd. Ed. New York: John
Wiley & Sons).
PLATINIC ACID, HEXACHLORO
[16941-12-1]
Formula: H2PtCl6; MW 409.81; crystallized as a hexahydrate, H2PtCl6 • 6H2O
Synonyms: chloroplatinic acid; hexachloroplatinic acid; hexachloro platinic(IV) acid
Uses
Chloroplatinic acid is used in preparing most platinum salts and complexes. It also is used as an electroplating bath for plating and coating of platinum. Other applications are in catalysis.
Physical Properties
The hexahydrate consists of red to brownish-red cubic crystals; deliquesces;
density 2.431g/cm3; melts at 60°C; very soluble in water and alcohol; soluble
in ether.
Preparation
Hexachloroplatinic acid is obtained in an intermediate step during extraction of platinum from minerals. The compound is formed when platinum is
dissolved in aqua regia containing a higher proportion of HCl and subsequently is evaporated repeatedly with hydrochloric acid, preferably in a chlorine atmosphere. Alternatively, hexachloroplatinic acid may be obtained by
dissolving platinum tetrachloride, PtCl4, in water.
Pure hexachloroplatinic acid may be prepared by dissolving platinum
sponge in hydrochloric acid under chlorine.
Reactions
Hexachloroplatinic acid decomposes completely when ignited, leaving a
residue of spongy platinum. Hexachloroplatinic acid on heating at 300°C in
chlorine forms platinum tetrachloride:
H2PtCl6 → PtCl4 + 2HCl
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PLATINUM
719
Reactions with ammonium chloride or other ammonium salts form a lemon
yellow precipitate of ammonium hexachloroplatinate, (NH4)2PtCl6 :
H2PtCl6 + 2NH4Cl → (NH4)2PtCl6 + 2HCl
Treatment with caustic alkali yields a white precipitate of hexahydroxoplatinic acid, H2Pt(OH)6:
H2PtCl6 + 6OH¯→ H2Pt(OH)6 + 6Cl¯
When hydrogen sulfide is bubbled into a boiling solution of hexachloroplatinic acid, a black precipatate of platinum disulfide, PtS2 (soluble in aqua
regia) is obtained:
H2PtCl6 + 2H2S → PtS2 + 6HCl
The above reaction is accompanied simultaneously with reduction of
[PtCl6]2— into platinum in metallic state. Metalic platinum, however, is a
minor product.
Addition of silver nitrate produces a yellow precipitate of silver hexachloroplatinate, Ag2PtCl6:
H2PtCl6 + 2AgNO3 → Ag2PtCl6 + 2HNO3
Hexachloroplatinic acid may be reduced to tetrachloroplatinic(II) acid,
H2PtCl4, by sulfur dioxide and other reducing agents.
Analysis
Elemental composition (anhydrous salt): Pt 47.60%, H 0.49%, Cl 51.90%.
The compound may be identified by its physical and chemical properties.
Platinum in an aqueous solution of the compound can be analyzed by flame
AA or ICP spectroscopy. Also, the compound can be measured by gravimetry
following precipitation with ammonium chloride, hydrogen sulfide, or silver
nitrate (see Reactions above).
PLATINUM
[7440-06-4]
Symbol Pt; atomic number 78; atomic weight 195.08; a Group VIII (Group 10)
noble metal; atomic radius 1.39Å; ionic radius of Pt2+ and Pt4+ in crystals having coordination numbers 4 and 6 are 0.60 Å and 0.63Å respectively; electron
configuration [Xe]4ƒ145d96s1; valence states +2, +3, +4, most common valence
+4; six stable isotopes: Pt-190 (0.011%), Pt-192 (0.80%), Pt-194 (32.96%), Pt195 (33.86%), Pt-196 (25.36%), Pt-198 (7.22%); twenty-nine radioactive isotopes in the mass range 168−190, 193, 197, 199−202; the longest lived radioactive isotope is naturally occurring Pt-190, t1/2 6.5x1011 years.
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PLATINUM
History, Occurrence, and Uses
Platinum was discovered in Colombia, South America by Ulloa in 1735 and
six years later in 1741 by Wood. The metal was isolated from native platinum
by de l’Isle in 1775 and produced in malleable form by Chabaneau in 1786.
Wollaston in 1803 developed a method of obtaining pure malleable platinum
from crude platinum by extraction with aqua regia. The process led to the discovery of two other platinum group metals, palladium and rhodium, that were
found in the aqua regia extract after platinum precipitated. Platinum derived
its name from platina originating from the Spanish word plata for silver,
because it was thought to be a trivial unwanted material associated with gold
in gold mines of Central America.
Platinum occurs in nature as a bright-white cubic crystalline solid with
metallic luster associated with other noble metals of its group. Platinum also
occurs as the mineral sperrylite, PtAs2, found as tin-white brittle cubic crystals containing 52−57% platinum in certain nickel-bearing deposits. Some
other minerals of platinum are cooperite PtS (Pt 80-86%); and braggite(Pt, Pd,
Ni)S (Pt 58-60%). The abundance of platinum in the earth’s crust is estimated to be 0.005 mg/kg.
Platinum metal and its alloys have numerous applications. As a precious
metal it is used extensively in jewelry. Other important applications include
construction of laboratory crucibles and high temperature electric furnaces; in
instruments as thermocouple elements; as wire; for electrical contacts; as
electrodes; in dentistry; in cigarette lighters; and for coating missile and jet
engine parts.
Platinum also is used extensively as a catalyst in hydrogenation, dehydrogenation, oxidation, isomerization, carbonylation, and hydrocracking. Also, it
is used in organic synthesis and petroleum refining. Like palladium, platinum
also exhibits remarkable ability to absorb hydrogen. An important application
of platinum is in the catalytic oxidation of ammonia in Ostwald’s process in
the manufacture of nitric acid. Platinum is installed in the catalytic converters in automobile engines for pollution control.
Physical Properties
Silvery-white lustrous metal; remains bright at all temperatures; face-centered cubic crystal; density 21.5g/cm3; Vickers hardness, annealed 38-40;
melts at 1,768.4°C; vaporizes at 3,825°C; vapor pressure at melting point
0.00014 torr; electrical resistivity 9.85 microhm-cm at 0°C; magnetic susceptibility 9.0x10—6 cm3/g; Poisson’s ratio 0.39; thermal neutron cross section 8
barns; insoluble in water and acids; soluble in aqua regia
Thermochemical Properties
∆Hƒ° (cry)
∆Hƒ° (gas)
∆Gƒ° (cry)
∆Gƒ° (gas)
S° (cry)
S° (gas)
0.0
135.1 kcal/mol
0.0
124.4 kcal/mol
9.94 cal/deg mol
46.0 cal/deg mol
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PLATINUM
Cρ (cry)
Cρ (gas)
∆Hfus
Thermal conductivity
Coefficient of linear expansion, at 20°C
721
6.19 cal/deg mol
6.09 cal/deg mol
5.30 kcal/mol
71.1 W/(m.K)
9.1x10–6/°C
Reactions
At ordinary temperatures platinum is inert to practically all substances
except aqua regia and, to a small extent, chlorine water. The metal is not
attacked by strong acids except aqua regia. It dissolves in aqua regia forming
chloroplatinic acid, H2PtCl6.
Platinum reacts with oxygen only at elevated temperatures. Finely divided
metal forms platinum oxide, PtO, at about 500°C. When heated at 1,000°C in
air or oxygen, platinum loses weight probably due to the evaporation of the
thin layer of PtO2 from its surface.
Fused alkalies, particularly potassium and barium hydroxides, are corrosive to platinum. In the presence of oxygen or oxidizing agents this corrosive
action of fused alkalies increases. Also, cyanide and nitrates of alkali metals
in fused state are corrosive to platinum.
Platinum combines with dry chlorine above 250°C forming platinum dichloride, PtCl2. Reaction with fluorine occurs at dull red heat forming platinum
tetrafluoride, PtF4, as the major product, with small amounts of difluoride,
PtF2.
Platinum can be alloyed with many elements at elevated temperatures.
Such elements include other noble metals, as well as, cobalt, selenium, silicon,
and arsenic and nonmetals like carbon, phosphorus, and sulfur.
Platinum, like palladium, absorbs a large volume of hydrogen, particularly
when heated. Hydrogen also diffuses through hot platinum sheet.
Platinum retains hydrogen at ordinary temperature and gives off the gas
when heated in vacuum.
Production
Platinum metal concentrate obtained after the mineral is subjected to various mechanical processes including froth flotation and gravity separation is
treated with aqua regia. Gold, platinum and palladium dissolve in aqua regia
leaving behind other noble metals and silver in the insoluble residues. Gold is
precipitated from the aqua regia extract by treating the solution with dibutyl
carbitol. Alternatively, gold may be removed from the chloride solution by
reduction with sulfur dioxide or ferrous salt to yield metallic gold. The filtrate
solution contains platinum and palladium in the form of chloroplatinic and
chloropalladic acids, H2PtCl6 and H2PdCl4, respectively. Ammonium chloride
is added to this solution to precipitate ammonium chloroplatinate (NH4)2PtCl6
leaving palladium in solution. The precipitate obtained at this stage contains
trace impurities. Crude complex is refined in a series of steps to obtain purified metal. Such refining steps may include igniting the complex; dissolving
the impure platinum sponge in aqua regia; treatment with sodium chloride to
precipitate sodium platinum chloride, Na2PtCl6, and converting pure
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PLATINUM DICHLORIDE
Na2PtCl6 to ammonium platinum chloride (NH4)2PtCl6. The purified ammonium complex is then ignited to form platinum sponge.
Analysis
Platinum in metallic form is brought into aqueous phase by boiling with
aqua regia and evaporating almost to dryness. This is followed by adding concentrated HCl and a small amount of NaCl and again evaporating to dryness.
Finally, the residue is dissolved in dilute HCl and diluted further for analysis. The aqueous solution is analyzed by flame-atomic absorption spectrophotometry using an air-acetylene flame. Measurement may be carried out at the
wavelength 265.9 nm. Platinum may be measured by other instrumental
techniques such as X-ray fluorescence and neutron activation analysis.
PLATINUM DICHLORIDE
[10025-65-7]
Formula: PtCl2; MW 265.99
Synonyms: platinum(II) chloride; platinous chloride
Uses
The compound does not have any notable commercial applications. It is
used to prepare tetrachloroplatinic(II) acid (choroplatinous acid) and tetrachloroplatinate salts.
Physical Properties
Olive green hexagonal crystals; density 6.05 g/cm3; decomposes to platinum
metal and chlorine on heating at 581°C; insoluble in water and alcohol; soluble in hydrochloric acid and ammonia solution.
Thermochemical Properties
∆Hƒ°
–29.5 kcal/mol
Preparation
Platinum dichloride is prepared by heating platinum sponge in chlorine at
about 500°C:
Pt + Cl2 → PtCl2
It also may be obtained by thermal decomposition of platinum tetrachloride, PtCl4 , or hexachloroplatinic acid:
PtCl4 → PtCl2 + Cl2
H2PtCl6 → PtCl2 + 2HCl +Cl2
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PLATINUM DIOXIDE
723
Reactions
Platinum dichloride dissolves in hydrochloric acid to form a dark brown
complex acid, tetrachloroplatinic(II) acid, H2PtCl4 in the solution:
PtCl2 + 2HCl → H2PtCl4
Tetrachloroplatinic(II) acid formed above may decompose to a small extent
forming metallic platinum and hexachloroplatinic(II) acid.
Reactions with carbon monoxide at moderate temperatures yield complexes [PtCl2(CO)]2, [PtCl2(CO)2], and [(PtCl2)2(CO)2], having melting points 194°,
142°, and 130°C, respectively.
Platinum dichloride forms complexes with ammonia, [Pt(NH3)4]Cl2, which
on heating yields [PtCl2(NH3)2].
Analysis
Elemental composition: Pt 73.36%, Cl 26.64%. The compound is dissolved
in concentrated HCl, diluted, and analyzed for platinum by flame-AA spectrophotometry (see Platinum). The salt may be identified by its olive green
color and other physical and x–ray properties. It forms a dark brown color in
HCl.
PLATINUM DIOXIDE
[1314-15-4]
Formula: PtO2; MW 227.08; forms mono-, di-, and tetrahydrates
Synonyms: platinum(IV) oxide; platinic oxide; Adams’ catalyst
Uses
Platinum dioxide, also known as Adams’ catalyst, is used commercially in
many hydrogenation reactions at ordinary temperatures, such as reduction of
olefinic and acetylenic unsaturation, aromatics, nitro, and carbonyl groups.
Physical Properties
Black solid; density 10.2 g/cm3; melts at 450°C; thermally decomposes;
insoluble in water, alcohol, acids and aqua regia; soluble in caustic potash
solution.
Thermochemical Properties
∆Hƒ° (g)
∆Gƒ° (g)
41.0 kcal/mol
40.1 kcal/mol
Preparation
Platinum dioxide is obtained as its monohydrate, PtO2•H2O, a brown-red
precipitate, upon boiling a solution of platinum tetrachloride, PtCl4, with sodium carbonate.
The anhydrous black dioxide, PtO2, may be prepared by treating a solution
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PLATINUM HEXAFLUORIDE
of hexachloroplatinic acid, H2PtCl6, with sodium carbonate. The yellow hexahydroxoplatinic acid, H2Pt(OH)6, is carefully heated below 100°C to yield the
black PtO2. Strong heating may decompose the dioxide to platinum metal.
Analysis
Elemental composition: Pt 85.91%, O 14.09%. The oxide may be characterized by its physical properties and by x-ray diffraction. The compound may be
thermally decomposed at elevated temperatures or reduced by hydrogen to
form platinum metal which may be digested with aqua regia and HCl, diluted, and analyzed by flame AA, ICP/AES or ICP/MS.
PLATINUM HEXAFLUORIDE
[13693-05-5]
Formula PtF6; ΜW 309.07; monomeric in vapor phase; Pt–F bond length is
about 1.82Å
Uses
Platinum hexafluoride does not have many commercial applications. It is
used as a strong oxidizing agent and can oxidize oxygen from the air. It is used
in research. Platinum hexafluoride forms compounds with molecular oxygen
and xenon, [O2+][PtF6–] and XePtF6 , respectively.
Physical Properties
Dark-red octahedral crystals; volatile and unstable; density 3.83g/cm3;
melts at 61.3°C; vaporizes at 69.14°C; reacts violently with water.
Preparation
Platinum hexafluoride may be prepared by heating platinum with fluorine
under pressure. The preparation should be in nickel or Monel apparatus as
the compound reacts with glass.
Reaction
The hexafluoride is a very powerful oxidizing agent reacting violently with
most oxidizable substances. Reaction with liquid water is violent forming HF,
oxygen, lower fluorides of platinum, and other products. In vapor phase
hydrolysis occurs more smoothly.
The hexafluoride decomposes on heating; also decomposed by UV radiation
to lower fluorides; and reacts with the inert gas xenon, forming a solid product, Xe(PtF6). It reacts with molecular oxygen to produce O2+PtF6– The compound attacks glass at ordinary temperatures.
Hazard
Platinum hexafluoride is dangerously corrosive. Inhalation of its vapors or
skin contact causes serious injury. Also, it can react explosively with a number of substances.
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PLATINUM MONOXIDE / PLATINUM TETRACHLORIDE 725
PLATINUM MONOXIDE
[12035-82-4]
Formula PtO; MW 211.08
Synonyms: platinum oxide; platinum(II) oxide
Uses
Platinum monoxide is used to prepare platinum-based catalysts.
Physical Properties
Violet-black solid; density 14.9g/cm3; decomposes on heating at 550°C;
insoluble in water and alcohol; soluble in aqua regia.
Preparation
Platinum monoxide is prepared by thermal decomposition of platinum(II)
hydroxide, Pt(OH)2, under careful heating.
Pt(OH)2 → PtO + H2O
If the hydroxide is heated too strongly and rapidly it disproportionates forming platinum metal and platinum dioxide:
2Pt(OH)2 → PtO2 + Pt + 2H2O
Platinum monoxide may be obtained as a black precipitate when an alkali
hydroxide is added to an aqueous solution of potassium tetrachloroplatinate(II) (potassium chloroplatinate), K2PtCl4.
Analysis
Elemental composition: Pt 92.41%, O 7.59%. The oxide can be identified by
its physical and x-ray properties. Additionally, platinum may be measured by
flame-AA following digestion of the solid with aqua regia and HCl (see
Platinum).
PLATINUM TETRACHLORIDE
[37773-49-2]
Formula: PtCl4; MW 336.89; also forms a pentahydrate, PtCl4 • 5H2O
Synonyms: platinum(IV) chloride; platinic chloride
Uses
Platinum tetrachloride is used to prepare chloroplatinic acid and many
platinum complexes, particularly with ammonia. Such complexes were pre-
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PLUTONIUM
pared and studied by Alfred Werner to support his theory on coordination
compounds.
Physical Properties
Brown-red crystalline solid; density 4.303g/cm3; decomposes at 370°C;
readily dissolves in water; dissolves in hydrochloric acid forming chloroplatinic acid, H2PtCl6; soluble in acetone; slightly soluble in ethanol; insoluble in
ether.
The pentahydrate PtCl4•5H2O constitutes red monoclinic crystals; density
2.43g/cm3; loses water on heating; very soluble in water; soluble in alcohol and
ether.
Thermochemical Properties
∆Hƒ° (g)
–55.4 kcal/mol
Preparation
Platinum tetrachloride is prepared by decomposition of hexachloroplatinic(IV) acid, H2PtCl6, in a stream of chlorine gas at 300°C.
Analysis
Elemental composition: Pt 52.56%, Cl 47.44%. Platinum tetrachloride may
be dissolved in water and analyzed for platinum (see Platinum). Also, it may
be identified by its physical properties and certain precipitation reactions
after dissolving in HCl (see Platinic Acid, Hexachloro).
PLUTONIUM
[7440-07-5]
Symbol Pu; atomic number 94; atomic weight 244; an actinide series
transuranium element; a man-made radioactive element; electron configuration [Rn]5ƒ67s2; partially filled ƒ subshell; valence states +3, +4, +5, +6; eighteen isotopes in the mass range 228-230, 232-246; all isotopes radioactive; the
longest lived isotope Pu-244, t1/2 8.2x107 year; the shortest lived isotope Pu233, t1/2 20.9 minute.
History, Occurrence, and Uses
Plutonium was discovered by Wahl, Seaborg, and Kennedy in 1941 at
Berkeley, California when they separated and identified its isotope of mass
238 produced from bombarding uranium isotopes with neutrons in a
cyclotron. In the same year the isotope Pu-239 was found to be fissionable.
However, only microgram quantities of Pu-239 were generated by cyclotron
bombardment. In 1943 Enrico Fermi and his group developed a process for
successful generation of much larger quantities of plutonium for nuclear
weapons. They achieved a self-sustaining nuclear chain reaction in a reactor
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PLUTONIUM
727
using uranium and graphite. This work eventually led to the first successful
testing of an atom bomb in the desert of New Mexico in July 1945.
Plutonium is the second transuranium element after neptunium. The element was named after the planet Pluto.
Plutonium is the most important transuranium element. Its two isotopes
Pu-238 and Pu-239 have the widest applications among all plutonium isotopes. Plutonium-239 is the fuel for nuclear weapons. The detonation power of
1 kg of plutonium-239 is about 20,000 tons of chemical explosive. The critical
mass for its fission is only a few pounds for a solid block depending on the
shape of the mass and its proximity to neutron absorbing or reflecting substances. This critical mass is much lower for plutonium in aqueous solution.
Also, it is used in nuclear power reactors to generate electricity. The energy
output of 1 kg of plutonium is about 22 million kilowatt hours. Plutonium-238
has been used to generate power to run seismic and other lunar surface equipment. It also is used in radionuclide batteries for pacemakers and in various
thermoelectric devices.
Physical Properties
Silvery-white metal; warm to touch because of its ionizing radiation; when
in appreciable amounts the metal can generate enough heat to boil water;
attains yellowish appearance when slightly oxidized. Six allotropic modifications are known: (1) alpha monoclinic form with sixteen atoms per unit cell;
stable at ordinary temperatures; density 19.86g/cm3; converts to beta form at
115°C. (2) beta form; body-centered monoclinic crystal structure; thirty-four
atoms per unit cell; density 17.70 g/cm3; stable between 115 to 200°C; converts
to gamma form at about 200°C. (3) gamma modification; face-centered
orthorhombic structure; eight atoms in unit cell; density 17.14g/cm3; exists
between 200 to 310°C; converts to delta form at 310°C. (4) delta allotrope;
face-centered cubic structure; four atoms per unit cell; density 15.92g/cm3;
stable in the temperature range 310 to 452°C; converts to a delta-prime form
at 452°C. (5) delta-prime form; body-centered tetragonal crystals; two atoms
per unit cell; density 16.00g/cm3; stable between 452 to 480°; converts to
another allotropic form, known as epsilon at 480°C. (6) epsilon form; body-centered cubic structure; two atoms per unit cell; density 16.51g/cm3; stable at
temperatures between 480 to 640°C.
Plutonium melts at 640°C; vaporizes at 3,228°C; electrical resistivity 146.4
microhm-cm at 0°C; Young’s modulus 14x106 psi; Poisson’s ratio 0.17; dissolves in concentrated hydrochloric, hydriodic, and perchloric acids (with
reaction).
Thermochemical Properties
∆Ηƒ°
Thermal conductivity
Coefficient of linear expansion
Cρ
Cρ (liquid at 675°C)
∆Ηα→β
0.0
0.0674 W/cmK
46.7 x 10-6/°C
8.84 cal/g-atom
10.0 cal/g-atom
900 ± 20 cal/g-atom
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PLUTONIUM
∆Ηβ→γ
∆Ηγ→δ
∆Ηδ→δ‘
∆Ηδ’→ε
∆Ηε→liquid
160 ± 10 cal/g-atom
148 ± 15 cal/g-atom
10 ± 10 cal/g-atom
444 ± 10 cal/g-atom
676 ± 10 cal/g-atom
Production
Plutonium is produced from natural uranium which is a mixture of nonfissionable uranium-238 (99.3%) and fissionable uranium-235(0.7%). The first
synthesis of this element was in a cyclotron generating plutonium in microgram quantities. The isotope Pu-239 can be produced in much larger quantities in a nuclear reactor, either a conventional thermal reactor or a breeder
type reactor by neutron bοmbardment of uranium- 238. The nuclear reactions
are shown below.
U + 01n→ 239
92 U + γ
238
92
U + 01n→ 239
92 U + γ
238
92
−
β
239
U →
93 Np
239
92
239
93
−
β
239
Np →
94 Pu
Higher isotopes such as Pu-240, -241, -242, etc. can be obtained from Pu239 by continued neutron bombardment.
Plutonium-239 also is produced from natural uranium by the so-called “pile
reactions” in which irradiation of uranium-235 isotope with neutrons produces fission, generating more neutrons and high energy (~200 MeV). These
neutrons are captured by the uranium-238 to yield plutonium-239.
Synthesis of plutonium in significant quantities requires a sufficiently long
reactor fuel irradiation period. Uranium, plutonium, and the fission products
obtained after neutron irradiation are removed from the reactor and stored
under water for several weeks. During such cooling periods most neptunium239 initially formed from uranium and present in the mixture transforms to
plutonium-239. Also, the highly radioactive fission products, such as xenon133 and iodine-131 continue to decay during this period.
Plutonium is recovered from uranium and fission products by solvent
extraction, precipitation, and other chemical methods. In most chemical
processes, plutonium first is converted to one of its salts, usually plutonium
fluoride, before it is recovered in purified metallic form. The fluoride is
reduced with calcium metal to yield plutonium. Electrorefining may produce
material of higher purity.
Plutonium is cast into small ingots by arc melting. All melting operations
must be carried out in vacuum or in an inert atmosphere to prevent any air
oxidation at high temperatures. Also, being a reactive metal, its recovery and
purification should be done in crucibles made of highly refractory and stable
materials.
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POLONIUM
729
Reactions
Plutonium is a reactive metal forming mostly tri-, tetra-, and hexavalent
compounds. The solutions of Pu3+ are blue. The trivalent Pu3+ is stable in solution in the absence of air. In the presence of air or oxygen, Pu3+ slowly oxidizes
to Pu4+. In cold acid medium, permanganate ion oxidizes Pu3+ to Pu4+. In
aqueous solutions Pu4+ salts impart pink or greenish color to the solutions.
Tetravalent Pu4+ converts to hexavalent plutonium, Pu6+ by the action of
–
strong oxidizing agents, such as dichromate, Cr2O72–, permanganate, MnO4
4+
or Ce salts.
The metal ion in higher oxidation states can be reduced by most common reducing agents, such as, sulfur dioxide, carbon monoxide, ferrocyanide
ion, hydrazine hydrochloride, and hydroxylamine hydrochloride to form Pu3+
(or Pu4+) ions in solution.
Plutonium combines with oxygen at high temperatures to form plutonium dioxide, PuO2 , and other oxides. The dioxide also is formed in the presence of water vapor. Ignition of the metal in air at 1,000°C yields PuO2.
Plutonium reacts with hydrogen at high temperatures forming
hydrides. With nitrogen, it forms nitrides, and with halogens, various plutonium halides form. Halide products also are obtained with halogen acids.
Reactions with carbon monoxide yields plutonium carbides, while with carbon
dioxide, the products are both carbides and oxides. Such reactions occur only
at high temperatures.
Plutonium forms several complexes in oxidation states +3, +4, and +6.
Hazard
Plutonium is one of the most dangerous substances known. The metal and
it’s salts are all highly toxic. Its ionizing radiation can cause cancer. The metal
can incorporate with bone marrow forming insoluble plutonium (IV) phosphate. The metal only leaves the body very slowly. All operations must be carried out by remote control devices with proper shields. In production, processing, handling, and storage of large quantities of plutonium or its compounds one must bear in mind its critical mass, which can vary with the shape
and the specific solid form or the quantities of plutonium contained in solutions.
POLONIUM
[7440-08-6]
Symbol Po; atomic number 84; atomic weight 209; a Group VIA (Group 16)
radioactive element; electron configuration [Xe]4ƒ145d106s26p4; valence states
–2, 0, +2, +4, +6; atomic radius 1.64Å; atomic volume 23.53cc/g-atom; the last
radioactive member of radium series; twentyfive isotopes; all radioactive; the
longest lived isotope is the alpha emitter Po–209, t1/2 105 ± 5 year.
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POLONIUM
History, Occurrence, and Uses
Polonium was discovered by Marie Curie in 1898 while investigating the
radioactivity of pitchblende. Mme. Curie named this new element after her
native country Poland. Polonium is a very rare element, found in exceedingly
small quantities in uranium ores. Its abundance in uranium ore is about
100mg/ton. Its applications are only a few. Polonium is used on brushes to
remove dusts from photographic film. It also is used in instruments to eliminate static charges. Polonium is used as a small source to generate alpha particles and neutrons; as a power source in devices where its radioactive decay
energy is converted into electrical energy. For this application the metal is
combined with lighter elements.
Physical Properties
Two crystalline forms exist; (1) alpha allotrope; a simple cubic low temperature form; density 9.196 g/cm3, and (2) beta modification: a rhombohedral
high temperature form; density 9.398 g/cm3
Both allotropic forms coexist between 18 to 54°C; melt at 254°C; vaporize
at 962°C; electrical resistivity 42 and 44 microhm-cm at 0°C for alpha- and
beta- forms, respectively; practically insoluble in water; soluble in dilute mineral acids.
Thermochemical Properties
∆Hƒ°
∆Gƒ°
∆Ηvap
∆Hsub
Coeff. linear expansion
0.0
0.0
24.6 kcal/mol
34.5 kcal/g-atom
(23.0±1.5)x10−6/°C
Production
Polonium can be recovered from natural pitchblende. The yield, however, is exceedingly small as 1 g of polonium is contained in about 25,000 tons
of pitchblende. The element may be isolated from the pitchblende extract by
deposition on a bismuth plate immersed in chloride solution.
Polonium can be produced from other sources, too, that offer much higher
yield than pitchblende. Two such processes are as follows:
(1) The element may be obtained from radioactive lead-210 (also,
known as RaD, the lead fraction in the extraction of radium from uranium
ore) by successive beta decay:
210
82
Pb
(RaD)
β–
−−−−−→
t11/2 22 yr
210
83
Bi
(RaE)
β–
−−−−−→
t11/2 22 yr
210
84
Po
(RaF)
The alpha emitter radioactive Po-210 that has a half-life of 138 days transforms to nonradioactive lead-206, the stable end product:
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POLONIUM
210
84
α−emission
Po t−−−−−−→
138 days
(RaF)
1/2
206
82
731
Pb(stable)
(RaG)
(2) Polonium also can be synthesized by neutron irradiation of natural bismuth in a reactor:
209
83
Bi +
1
0
n
β–
210
→ 210
83 Bi −−→ 83 Po
After neutron irradiation bismuth (canned in aluminum jackets) is dis–
solved in a mixture of hydrochloric and nitric acids and excess NO3 is
removed by adding a reducing agent, such as, urea or formic acid. If bismuth
is used as an anode, the reducing agent is dissolved in HCl. Various methods
are applied for concentration of polonium in the acid mixture and its subsequent separation from bismuth. Such processes include spontaneous deposition of polonium over a less electropositive metal and coprecipitation with tellurium. In the latter method, a Te4+ or Te6+ salt is added to the extract, followed by addition of stannous chloride, which reduces both the tellurium and
polonium to their metallic state, coprecipitating them from bismuth in the
extract mixture.
Another method to separate polonium from bismuth involves heating at
650°C to convert the metals into their oxides. This is followed by further heating to about 800°C at reduced pressure in which polonium metal is removed
by volatilization.
Polonium may be purified by various processes. Such purification methods
include precipitation of polonium as sulfide and then decomposing the sulfide
at elevated temperatures; spontaneous decomposition of polonium onto a
nickel or copper surface; and electrolysis of nitric acid solutions of poloniumbismuth mixture. In electrolytic purification polonium is electrodeposited onto
a platinum, gold, nickel, or carbon electrode.
Reactions
Polonium resembles tellurium, the element above it in the same Group, in
chemical behavior.
At ordinary temperatures polonium oxidizes slowly in air forming the basic
oxide, PoO2:
Po + O2 → PoO2
The metal dissolves in dilute hydrochloric acid forming pink-red polonium
dichloride:
Po + 2HCl → PoCl2 + H2
The unstable dichloride converts to yellow tetrachloride, PoCl4.
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POTASSIUM
Polonium dissolves in concentrated nitric acid and aqua regia, oxidizing to
Po4+ state. Reaction with nitric acid forms adducts that probably have the
compositions 4PoO2•N2O5; 4PoO2•3N2O5 and Po(NO3)4•N2O4. The metal also
dissolves in concentrated sulfuric and selenic acids forming polonium sulfate,
Po(SO4)2 and Po(SeO4)2 , respectively. Another product, 2PoO2•SO3, also has
been identified.
Because of its radioactivity and alpha emission, polonium forms many
types of radiolytic oxidation-reduction products.
Analysis
At trace levels, polonium can be separated effectively by solvent extraction,
ion exchange, paper chromatography, and other techniques. Diisopropyl
ketone, di-n-octylamine, and tri-n-butylphosphate are suitable solvents for
extraction. Trace amounts of polonium in solutions or solid mixtures containing no other emitters can be determined by measuring its alpha activity.
Hazard
As with other radioactive substances, exposure to its ionizing radiation can
cause cancer. When ingested it tends to accumulate in the liver, kidney, and
spleen causing radiation damage from the alpha particles. All operations and
handling must be carried out in leak-proof boxes by mechanical means behind
thick neutron shields.
POTASSIUM
[7440-09-7]
Symbol K; atomic number 19; atomic weight 39.098; a Group 1A (Group1)
alkali metal element; atomic radius 2.35Å; ionic radius, K+ 1.33Å; electron
configuration [Ar]4s1; valence state +1; ionization potential 4.341eV; standard
redox potential, Eº K+ + e− ↔ K(s) –2.925V; three natural isotopes: K39(93.258%), K-40 (0.0117%) and K-41 (6.730%); naturally occurring K-40 is
radioactive, t1/2 1.25x109 year, beta emitter; fourteen synthetic radioisotopes
in the mass range 35–38 and 42–51.
History, Occurrence, and Uses
Potassium was first isolated as a free metal in 1807 by Sir Humphry Davy.
It was the first alkali metal to be discovered, produced by electrolysis of potassium carbonate (potash). The element was earlier called Kalium, derived from
the Arabic word qili, meaning grass wort, the ash of which was a source of
potash. The element derived its symbol K from Kalium. The English name
potassium came from potash (pot ash), the carbonate salt of the metal.
Potassium is distributed widely in nature. The metal is too reactive to occur
in native elemental form. It is the seventh most abundant element on earth,
constituting 2.40% by weight of the earth’s crust. It is abundantly present in
sea water. Oceans contain 0.07% (wt to volume) potassium chloride.
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POTASSIUM
733
Potassium occurs in many igneous rocks, such as, feldspar (potassium aluminum silicate), KAlSi3O8 (leucite) and mica, KH2Al3(SiO4)3. Disintegration
of these rocks adds potassium to soil and water. Deposits of potassium chloride are found in practically all salt beds, associated with sodium chloride.
Some important potassium minerals are leucite, KAlSi2O6; glauconite (a complex silicoaluminate structure of varying compositions); sylvite, KCl; carnallite, KCl•MgCl2•6H2O; langbeinite, K2SO4•2MgSO4; and polyhalite,
K2SO4•MgSO4•2CaSO4•2H2O.
Potassium, along with nitrogen and phosphorus, is an essential element
needed for plant growth. In plants, it occurs mostly as K+ ion in cell juice. It
is found in fruit or seed. Deficiency can cause curling leaves, yellow or brown
coloration of leaves, weak stalk and diminished root growth. Potassium deficiency has been associated with several common animal ailments. Potassium
is in extracellular fluid in animals at lower concentrations than sodium.
Physical Properties
Silvery metal; body-centered cubic structure; imparts crimson-red color to
flame; density 0.862g/cm3 at 20ºC; melts at 63.25ºC; density of liquid potassium at 100ºC is 0.819 g/cm3 and 0.771g/cm3 at 300ºC; vaporizes at 760ºC; vapor
pressure 123 torr at 587ºC; electrical resistivity 6.1 microhm-cm at 0ºC and
15.31 microhm-cm at 100ºC; viscosity 0.25 centipoise at 250ºC; surface tension
86 dynes/cm at 100ºC; thermal neutron absorption cross section 2.07 barns;
reacts violently with water and acids; reacts with alcohol; dissolves in liquid
ammonia and mercury
Thermochemical Properties
∆Hƒ° (cry)
∆Hƒ° (gas)
∆Gƒ° (cry)
∆Gƒ° (gas)
S°(cry)
S°(gas)
Cρ (cry)
Cρ (gas)
∆Hfus
∆Hvap
Thermal conductivity at 200ºC
0.0
21.33 kcal/mol
0.0
14.49 kcal/mol
15.34 cal/deg mol
38.29 cal/deg mol
7.07 cal/deg mol
4.97 cal/deg mol
0.555 kcal/mol
0.496 kcal/mol
44.77 W/m.K
Production
Potassium can be produced by several methods that may be classified
under three distinct types: (1) electrolysis, (2) chemical reduction, and (3)
thermal decomposition.
Electrolysis processes have been known since Davy first isolated the metal
in 1807. Electrolysis, however, suffers from certain disadvantages. A major
problem involves miscibility of the metal with its fused salts. Because of this
molten potassium chloride, unlike sodium chloride, cannot be used to produce
the metal. Fused mixtures of potassium hydroxide and potassium carbonate
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POTASSIUM
or chloride have been used as electrolytes with limited success.
Chemical reduction processes are employed nowadays in commercial, as
well as, laboratory preparation of potassium. In one such process, molten
potassium chloride is reduced with sodium at 760 to 880ºC and the free metal
is separated by fractionation:
KCl + Na → K + NaCl
Potassium is obtained at over 99.5% purity. The metal, alternatively, may
be alloyed with sodium for further applications.
Reduction of potassium fluoride with calcium carbide at 1,000 to 1,100ºC
(Greisheim process) is an effective production method (Greer, J.S., Madaus,
J.H and J.W. Mausteller. 1982. In Kirk-Othmer Encyclopedia of Chemical
Technology, 3rd ed. p. 914, New York: Wiley Interscience):
2KF + CaC2 → CaF2 + 2C + 2K
Some other chemical reduction methods that may be applied for laboratory
generation of small quantities of potassium from its salts at high temperatures require a suitable reducing agent such as carbon, calcium, or calcium
carbide:
K2CO3 + 2C → 3CO +2K
2KCl + Ca → CaCl2 + 2K
2KCl + CaC2 → CaCl2 + 2C + 2K
2K2CO3 + +3Si + 3CaO → 4K + 2C + 3CaSiO3
2K2SiO3 + Si + 3 CaO → 4K + 3CaSiO3
Potassium can be produced by thermal decomposition of potassium azide:
2KN3 → 2K + 3N2
High purity metal may be produced by distillation of technical grade metal.
Potassium (technical grade) may be packed under nitrogen. Argon should be
used for packing high purity metal. Metal is shipped in stainless steel or carbon containers. In small quantities potassium is transported in glass or metal
ampules.
Reactions
Potassium reacts with oxygen or air forming three oxides: potassium
monoxide, K2O; potassium peroxide, K2O2; and potassium superoxide, KO2.
The nature of the product depends on oxygen supply. In limited supply of oxygen potassium monoxide is formed, while in excess oxygen, superoxide is
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POTASSIUM
735
obtained:
4K + O2 → 2K2O
2K + O2 → K2O2
K + O2 → KO2
Potassium reacts violently with water, forming potassium hydroxide:
2K + 2H2O → 2KOH + H2
Potassium reacts with hydrogen at about 350ºC to form potassium hydride:
2K + H2 → 2KH
Reactions with halogens, fluorine, chlorine and bromine occur with explosive violence. Thus, in contact with liquid bromine it explodes forming potassium bromide:
2K + Br2 → 2KBr
Potassium ignites in iodine vapor forming potassium iodide.
Violent reactions can occur with many metal halides. For example, with
zinc halides or iron halides, single replacement reactions take place. Such
potassium-metal halide mixtures can react violently when subjected to
mechanical shock.
At ordinary temperatures, potassium does not combine with nitrogen but
with an electric charge, potassium azide is formed.
Reaction with carbon (graphite) at above 400ºC produces a series of
carbides, such as KC4, KC8, and KC24. With carbon monoxide, an unstable
explosive carbonyl forms:
K + CO → KCO
Potassium reduces carbon dioxide to carbon, carbon monoxide and
potassium carbonate:
6K + 5CO2 → CO + C + 3K2CO3
Potassium reacts with ammonia gas to form potassium amide with liberation
of hydrogen:
2K + 2NH3 → 2KNH2 + H2
Reactions with phosphorus, arsenic and antimony form phosphide,
arsenide, and antimonide of potassium, respectively:
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POTASSIUM ACETATE
K + As → K3As
Reaction with sulfur forms three sulfides. When reactants are in molten
state, the product is K2S, but in liquid ammonia K2S2 and KS2 are the main
products.
Potassium reacts explosively with sulfuric acid, forming potassium sulfate
with evolution of hydrogen:
K + H2SO4 → K2SO4 + H2
Potassium liberates hydrogen from ethanol forming potassium ethoxide:
2K + 2C2H5OH → 2C2H5OK + H2
Reaction with potassium nitrate yields potassium monoxide and nitrogen:
10K + 2KNO3 → 6K2O + N2
Analysis
Potassium and its salts can be identified by flame test. It imparts lilac color
to the flame. Potassium ion in aqueous solution can be identified by reaction
with sodium tetraphenylborate, NaB(C6H5)4. In weakly acid solution, a white
precipitate of the potassium salt KB(C6H5)4 is obtained. The precipitate is filtered, dried, and weighed to measure potassium. The test is quantitative.
Potassium at trace concentrations in aqueous samples can be measured by
a flame photometer at a wavelength of 766.5 nm. Either a flame photometer
or an atomic absorption spectrometer operating in flame emission mode can
be used for such analysis.
Potassium also can be measured by ICP/AES. The wavelengths at which it
can be analyzed without interference from other metals are 766.49 and 769.90
nm. Other wavelengths may be used. Potassium ion in aqueous solution can
be identified quantitatively by using a potassium ion-selective electrode
attached to a pH meter having an expanded millivolt scale or to a specific ion
meter having a direct readout concentration scale for potassium.
Hazard
Potassium metal can be dangerous to handle if proper precautions are not
taken. Many of its reactions at ordinary temperatures can proceed to explosive violence (see Reactions). Also, it liberates flammable hydrogen gas when
combined with water, acids, and alcohols.
POTASSIUM ACETATE
[127-08-2]
Formula: KC2H3O2 or CH3COOK; MW 98.14
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POTASSIUM BICARBONATE 737
Uses
Potassium acetate is used in the manufacture of glass; as a softening agent
for papers and textiles; as a dehydrating agent; and as a buffer. In medicine
it is used as an expectorant and diuretic.
Physical Properties
White lustrous powder or colorless deliquescent crystals; density 1.57
g/cm3; melts at 292ºC; highly soluble in water, 253g/100mL at 20ºC, more soluble in hot water, 492g/100mL at 62ºC; aqueous solution alkaline, pH of 0.1M
solution 9.7; soluble in methanol, ethanol and liquid ammonia; insoluble in
ether and acetone.
Thermochemical Properties
∆Ηƒ° (cry)
∆Ηƒ° (amp)
∆Gƒ° (amp)
S° (amp)
Cρ (amp)
–172.8 kcal/mol
–176.5 kcal/mol
–156.0 kcal/mol
45.2 cal/deg mol
3.7 cal/deg mol
Preparation
Potassium acetate is prepared by addition of potassium carbonate in a
small volume of water to acetic acid solution, followed by evaporation and
crystallization:
K2CO3 + 2CH3COOH → 2CH3COOK + H2O
Analysis
Elemental composition: K 39.85%, C 24.48%, H 3.08%, O 32.60%.
Potassium may be identified by flame testing. An aqueous solution can be
analyzed for potassium by flame photometry, ICP/AES, or ion selective electrode (see Potassium). Acetate anion may be measured in aqueous solution by
ion chromatography under appropriate conditions.
POTASSIUM BICARBONATE
[298–14–6]
Formula KHCO3; MW 100.12
Synonyms: potassium hydrogen carbonate; potassium acid carbonate
Uses
Potassium bicarbonate is used in baking powder and effervescent salts. In
medicine, the salt is a gastric antacid and an electrolyte replenisher. It also is
dry powder in fire extinguishers.
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POTASSIUM BISULFIDE
Physical Properties
Colorless transparent crystal or white powder; monoclinic structure; density 2.17 g/cm3; decomposes above 100°C; soluble in water, 22.49 g/100ml at
20°C, 60 g/100ml at 60°C pH of 0.1M aqueous solution 8.2; practically insoluble in alcohol.
Preparation
Potassium bicarbonate is obtained by passing carbon dioxide through a
cold, concentrated solution of potassium carbonate:
K2CO3 + CO2 + H2O → 2 KHCO3
Alternatively, KHCO3 is produced by passing excess carbon dioxide
through aqueous potassium hydroxide. At first, potassium carbonate is
formed which then converts to the bicarbonate as shown in the above reaction.
Potassium bicarbonate cannot be made by Solvay process because of
its high solubility in water.
Reactions
Heating the bicarbonate yields normal carbonate, liberating carbon dioxide
and water:
2KHCO3 → K2CO3 + CO2 ↑ + H2O↑
When the salt is added to dilute acids, carbon dioxide is liberated:
KHCO3 + HCl → K+ + Cl¯ + CO2↑ + H2O
Reaction with caustic potash in solution forms potassium carbonate:
KHCO3 + KOH → K2CO3 + CO2 + H2O
Analysis
Elemental composition: K 39.05%, C 11.99%, H 1.01%, O 47.94%.
Potassium may be analyzed by AA spectroscopy in emission mode or by flame
photometry (see Potassium). The aqueous solution may be treated with HCl
and the CO2 evolved may be noted from effervescence and tested by GC-TCD
or by GC/MS. The characteristic mass ion for CO2 is 44. Alternatively, the
HCO3– anion or the CO32– anion (converted by heating the bicarbonate) may
be identified by ion chromatography.
POTASSIUM BISULFIDE
[1310–61–8]
Formula KHS; MW 72.17; usually exists as a hemihydrate
Synonyms: potassium hydrosulfide; potassium hydrogen sulfide; potassium
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POTASSIUM BOROHYDRIDE
739
sulfhydrate
Preparation
Potassium bisulfide is made by reacting calcium hydrogen sulfide with
potassium sulfate:
Ca(HS)2 + K2SO4 → 2KHS + CaSO4
The compound can be made by reacting hydrogen sulfide with potassium
sulfide:
H2S + K2S → 2KHS
Purer compound may be produced by passing dry hydrogen sulfide through
a solution of potassium metal dissolved in absolute ethanol:
2H2S + 2K → 2KHS + H2
Physical Properties
Colorless crystals or white crystalline mass; rapidly deliquesces; converts
to a yellow rhombohedral crystalline mass upon exposure to air, forming polysulfides and H2S; density 1.68g/cm3; the hemihydrate loses water at about
175°C; melts at 455°C to a dark red liquid; decomposes in water; soluble in
alcohol.
Thermochemical Properties
∆Ηƒ°
∆Ηsoln (at 17°C)
∆Ηsoln (hemihyadrate at 16°C)
62.5 kcal/mol
0.77 kcal/mol
0.62 kcal/mol
Analysis
Elemental composition: K 54.18%, S 44.42%, H 1.40%. An aqueous solution
may be analyzed for potassium by various methods (see Potassium). The compound on exposure to air evolves H2S which can be detected from its odor, as
well as by various tests (see Hydrogen Sulfide).
POTASSIUM BOROHYDRIDE
[13762–51–1]
Formula KBH4; MW 53.95
Synonym: potassium tetrahydroborate
Uses
Potassium borohydride, unlike sodium borohydride, has very limited applications. The compound is a reducing agent.
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POTASSIUM BROMATE
Physical Properties
White crystalline solid; stable in air; nonhygroscopic; density
1.11g/cm3; decomposes at about 500°C without melting; soluble in water,
19g/100ml at 25°C; stable in alkaline solution; soluble in liquid ammonia and
dimethyl formamide; slightly soluble in methanol, 0.7 g/100ml at 20°C.
Preparation
Potassium borohydride may be prepared by reacting potassium hydroxide
with sodium borohydride. The salt precipitates from an aqueous solution of
sodium borohydride with addition of potassium hydroxide:
NaBH4 + KOH → KBH4 + Na+ + OH¯
Also, potassium borohydride can be made by reacting potassium hydride
with methyl borate at high temperature:
4 KH + B(OCH3)3 → KBH4 + 3KOCH3
Potassium borohydride also may be prepared by reacting potassium
tetramethoxyborohydride with diborane at low temperatures; or by passing
diborane through a solution of potassium methylate in methanol.
Analysis
Elemental composition: K 72.47%, B 20.06%, H 7.47%. The salt is dissolved
in water and the solution analyzed for potassium and boron (see Potassium
and Boron).
POTASSIUM BROMATE
[7758–01–2]
Formula: KBrO3; MW 167.00
Uses
Potassium bromate is an oxidizing reagent in bromate-bromide mixture for
titrimetric analysis. It also is a bread- and flour-improving agent.
Physical Properties
Colorless trigonal crystals or fine white crystals or granules; density 3.27
g/cm3 at 18°C; melts at 350°C; decomposes at about 370°C evolving oxygen;
moderately soluble in water, 13.3 g/100mL at 40°C; slightly soluble in alcohol;
insoluble in acetone.
Thermochemical Properties
∆Ηƒ°
∆Gƒ°
–86.10 kcal/mol
–64.82 kcal/mol
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POTASSIUM BROMIDE
S°
Cρ
741
35.65 cal/deg mol
28.72 cal/deg mol
Preparation
Potassium bromate can be produced by electrolysis of potassium bromide
solution. Alternatively, the compound is obtained by adding potassium bromide to a saturated solution of sodium bromate or calcium bromate. The salt
is recovered from solution by crystallization.
Analysis
Elemental composition: K 23.41%, Br 47.85%, O 28.74%. Aqueous solution
of the salt after sufficient dilution may be analyzed for its potassium content
by AA, ICP, or flame photometry (see Potassium) and for bromate anion by ion
chromatography. Also, bromate content can be measured by iodometric titration using a standard solution of sodium thiosulfate and starch as indicator.
The redox reactions are as follows:
BrO3¯ + 5 Br¯ (excess) + 6H+ → 3Br2 + 3H2O
Br2 + 2I¯ → 2Br¯ + I2
Liberated iodine is titrated against a standard solution of thiosulfate until the
starch solution’s blue decolorizes.
Toxicity
Ingestion of the salt or its solution can cause nausea, vomiting, diarrhea,
and renal injury. Also, it can induce methemoglobinemia.
POTASSIUM BROMIDE
[7758–02–3]
Formula: KBr; MW 119.00
Uses
Potassium bromide is used to make photographic plates and papers and in
engraving. Other uses are as a brominating agent in organic synthesis and in
the bromate-bromide mixture in titrimetric analysis. In medicine potassium
bromide is a sedative and anticonvulsant.
Physical Properties
Colorless cubic crystals or white granules or powder; density 2.75 g/cm3 at
25°C; melts at 734°C; vaporizes at 1,435°C; readily dissolves in water, solubility at 0°C 53.5 g/100mL and at 100°C 102 g/100mL; aqueous solution neutral; soluble in glycerol, 21.7 g/100mL; sparingly soluble in boiling ethanol
4.76 g/100mL.
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POTASSIUM BROMIDE
Thermochemical Properties
∆Ηƒ° (cry)
∆Ηƒ° (gas)
∆Gƒ° (cry)
∆Gƒ° (gas)
S° (cry)
S° (gas)
Cρ(cry)
Cρ(gas)
–94.12 kcal/mol
–43.04 kcal/mol
–90.98 kcal/mol
–50.89 kcal/mol
22.92 cal/deg mol
59.95 cal/deg mol
12.50 cal/deg mol
8.52 cal/deg mol
Preparation
Potassium bromide is prepared by reacting bromine with potassium carbonate:
3K2CO3 + 3Br2 → KBrO3 + 5KBr + 3CO2
Potassium bromate, KBrO3, is less soluble than the bromide. Thus, most
potassium bromate may be removed by filtration. Remaining bromate can be
converted to bromide by reduction with iron. After filtering iron from the solution, potassium bromide is obtained by evaporation and crystallizaton.
Another method of preparation involves treating bromine with warm concentrated aqueous solution of potassium hydroxide:
3Br2 + 6KOH → 5KBr + KBrO3 + 3H2O
Bromide-bromate solution is evaporated to dryness. The residue is heated
with charcoal:
2KBrO3 + 3C → 2KBr + 3CO2
Potassium bromide also can be prepared by treating iron turnings with a
35 wt% aqueous solution of bromine. The product ferrosoferric bromide is
boiled in potassium carbonate solution containing a slight excess of 15%
potassium carbonate (Dancy, W.B. 1980. Potassium Compounds. In KirkOthmer Encyclopedia of Chemical Technology, 3rd ed. p. 963. New York: Wiley
Interscience). The method does not involve bromate formation. The second
step of the process may be represented in the following reaction:
Fe3Br8•16H2O + 4K2CO3 → 8KBr + 4CO2 + Fe3O4 + 16H2O
Potassium bromide also can be produced by electrolytic process.
Analysis
Elemental composition: K 32.85%, Br 67.15%. Potassium can be determined in solid form by flame testing. In aqueous solution, potassium can be
measured by flame photometry, ICP/AES or electrode methods. Bromide ion
can be analyzed in aqueous solution by ion chromatography.
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POTASSIUM CARBONATE
743
Toxicity
Potassium bromide ingested in large doses can cause CNS depression.
Other symptoms of chronic intake are mental deterioration and an acne-type
skin eruption.
POTASSIUM CARBONATE
[584–08–7]
Formula: K2CO3; MW 138.21
Synonyms: potash; pearl ash; salt of tartar
Occurrence and Uses
Potassium carbonate occurs in wood ashes. It is one of the first known salts
of potassium and was used historically in recovering metalic potassium. The
compound has numerous potential applicatons. However, in most cases the
cheaper and equivalent sodium carbonate is used. An important application
of potassium carbonate involves making specialty television glass. Other
applicatons are in pottery; soaps and liquid shampoos; process engraving and
lithography; to depress the freezing point of water in fire extinguishers for
unheated warehouses; and in tanning and leather work. An important use of
this compound is preparing several other potassium salts.
Physical Properties
Colorless monoclinic crystals or granular powder; hygroscopic; density
2.428 g/cm3 at 20°C; melts at 891°C; decompses on further heating; very soluble in water 112 g/100mL at 20°C and more soluble in boiling water, 156
g/100mL at 100°C; aqueous solution strongly alkaline; insoluble in alcohol
and acetone.
Thermochemical Properties
∆Ηƒ°
∆Gƒ°
S°
Cρ
–275.1 kcal/mol
–254.2 kcal/mol
37.17 cal/deg mol
27.35 cal/deg mol
Preparation
Potassium carbonate is produced most conveniently by passing carbon
dioxide into an aqueous solution of caustic potash, evaporating the solution to
obtain the bicarbonate, and heating the bicarbonate:
KOH + CO2 → KHCO3
2KHCO3 → K2CO3 + CO2↑ + H2O↑
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POTASSIUM CARBONATE
The carbonate salt also can be prepared by heating potassium formate in
air or oxygen:
2HCOOK + O2 → K2CO3 + CO2↑ + H2O↑
Potassium formate obtained from purified producer gas (see Potassium
Formate) is heated in a rotary furnace having free access to air.
At ordinary temperatures, the carbonate salt crystallized from water is
obtained as a dihydrate, K2CO3•2H2O
The carbonate also can be made from potassium chloride, magnesium carbonate trihydrate and carbon dioxide under 30 atm at ordinary temperatures
by Engel-Precht process:
2KCl + 3 MgCO3•3H2O + CO2 → 2KHCO3•MgCO3•4H2O + MgCl2
The hydrated double salt on ignition decomposes giving potassium carbonate that may be extracted with water:
2KHCO3•MgCO3•4H2O → K2CO3 + 2MgCO3 + 9H2O + CO2
Small amounts of potassium carbonate were derived historically from
leaching wood ash. The process is now obsolete.
Reactions
When carbon dioxide is passed into an aqueous solution of potassium carbonate, potassium bicarbonate is produced:
K2CO3 + CO2 + H2O → 2KHCO3
Reactions with dilute acids evolve carbon dioxide:
K2CO3 + H2SO4 → 2K+ + SO42– + CO2↑+ H2O
Potassium carbonate-carbon mixture reacts with ammonia at high temperatures to form potassium cyanide:
K2CO3 + 4C + 2NH3 → 2KCN + 3CO↑ + 3H2 ↑
Analysis
Elemental composition: K 56.58%, C 8.69%, O 34.73%. The salt can be identified from its physical and chemical properties. Its aqueous solution is highly alkaline. Reaction with dilute acids evolves CO2 with effervescence. The latter can be identified by GC–TCD or GC/MS. The primary characteristic mass
ion for CO2 is 44. Also, CO3 2– anion can be measured by ion chromatography.
Potassium can be analyzed by various instrumental and wet methods (see
Potassium).